XI Chem Ch4 ChemicalBonding&MolecularStructure Concept

8/6/2019 XI Chem Ch4 ChemicalBonding&MolecularStructure Concept

1/12

Get the Power of Visual Impact on your sideLog on to www.topperlearning.com

1

Class XI: Chemistry

Chapter 4: Chemical Bonding and Molecular Structure

Top Concepts

1. The attractive force which holds together the constituent particles (atoms, ionsor molecules) in chemical species is known as chemical bond.

2. Tendency or urge atoms of various elements to attain stable configuration of

eight electrons in their valence shell is cause of chemical combination.

3. The principle of attaining a maximum of eight electrons in the valence shell or

outermost shell of atoms is known as octet rule.

4. Electronic Theory: Kossel-Lewis approach to chemical Bonding: Atoms achievestable octet when they are linked by chemical bonds. The atoms do so either

by transfer or sharing of valence electrons. Inner shell electrons are not

involved in combination process.

5. Lewis Symbols or electron dot symbols: The symbol of the element represents

the whole of the atom except the valence electrons (i.e. nucleus and the

electrons in the linear energy shells). The valence electrons are represented by

placing dots (.) or crosses (x) around the symbol.

6. Significance of Lewis Symbols: The Lewis symbols indicate the number of

electrons in the outermost or valence shell which helps to calculate common or

group valence.

7. The common valence of an element is either equal to number of dots or

valence electrons in the Lewis symbol or it is equal to 8 minus the number of

dots or valence electrons.

8. The bond formed by mutual sharing of electrons between the combining

atoms of the same or different elements is called a covalent bond.

9. If two atoms share one electron pair, bond is known as single covalent bond

and is represented by one dash ().


2/12


2

10. If two atoms share two electron pairs, bond is known as double covalent bond

and is represented by two dashes ( =).

11. If two atoms share three electron pairs, bond is known as triple covalent bond

and is represented by three dashes ( ).

12. The formal charge of an atom in a polyatomic ion or molecule is defined as the

difference between the number of valence electrons in an isolated (or free)

atom and the number of electrons assigned to that atom in a Lewis structure.

It may be expressed as:

Formal chargeon an atom =

in free atom

Number of Number of Number of 1

valence electrons nonbonding bonding2

in free atom (lone pair) electrons (shared) electrons

13. Significance of Formal charge: The formal charges help in selection of lowest

energy structure from a number of possible Lewis structures for a given

molecule or ion. Lowest energy structure is the one which has lowest formal

charges on the atoms.

14. Expanded octet: Compounds in which central atom has more than eight

electrons around it, atom is said to possess an expanded octet.

15. Exceptions to the Octet Rule:

Hydrogen molecule: Hydrogen has one electron in its first energy shell ( n =

1) It needs only one more electron to fill this shell, because the first shell

cannot have more than two electrons. This configuration (1s2

) is similar tothat of noble gas helium and is stable. In this case, therefore, octet is not

needed to achieve a stable configuration.

Incomplete octet of the central atom : The octet rule cannot explain the

formation of certain molecules of lithium, beryllium, boron, aluminum, etc.


3/12


3

(LiCl, BeH 2 , BeCl 2 , BH 3 , BF 3 ) in which the central atom has less than eight

electrons in the valence shell as shown below:

Expanded octet of the central atom : There are many stable molecules

which have more than eight electrons in their valence shells. For example,PF 5 , has ten; SF 6 has twelve and IF 7 ha fourteen electrons around the

central atoms, P, S, and I respectively.

Odd electron molecules: There are certain molecules which have odd

number of electrons, like nitric oxide, NO and Nitrogen dioxide, NO 2 . In

these cases, octet rule is not satisfied for all the atoms.

It may be noted that the octet rule is based upon the chemical inertness of

noble gases. However, it has been found that some noble gases (especially

xenon and krypton) also combine with oxygen and fluorine to form a large

number of compounds such a XeF 2 , KrF 2 , XeOF 2 , XeOF 4 , XeF 6 , etc.

This theory does not account for the shape of the molecules.

It cannot explain the relative stability of the molecule in terms of the

energy.

16. General Properties of Covalent Compounds:

1. The covalent compounds do not exist as ions but they exist as

molecules.

2. The melting and boiling points of covalent compounds are generally low.

3. Covalent compounds are generally insoluble or less soluble in water and

other polar solvents. However, these are soluble in non- polar solvents.

4. Since covalent compounds do not give ions in solution, these are poor

conductors of electricity in the fused or dissolved state.

5. Molecular reactions are quite slow because energy is required to break

covalent bonds.

6. Since the covalent bond is localized in between the nuclei of atoms, it is

directional in nature.


4/12


4

17. Co-Ordinate Covalent Bond:

Covalent type bond in which both the electrons in the shared pair come

from one atom is called a coordinate covalent bond.

Co- Ordinate Covalent Bond is usually represented by an arrow ( )pointing from donor to the acceptor atom.

Co- Ordinate Covalent bond is also called as dative bond, donor acceptor

bond, semi- polar bond or co-ionic bond.

18. The electrostatic force of attraction which holds the oppositely charged ions

together is known as ionic bond or electrovalent bond.

19. Ionic compounds will be formed more easily between the elements withcomparatively low ionization enthalpy and elements with comparatively high

negative value of electron gain enthalpy.

20. A quantitative measure of the stability of an ionic compound is provided by its

lattice enthalpy and not simply by achieving octet of electrons around the ionic

species in the gaseous state.

21. Lattice enthalpy may also be defined as the energy required to completely

separating one mole of a solid ionic compound into gaseous ionic constituents.

22. Factor affecting lattice enthalpy:

Size of the ions: Smaller the size of the ions, lesser is the inter-nuclear

distance and higher will be lattice enthalpy.

Larger the magnitude of charge on the ions, greater will be the attractive

forces between the ions. Consequently, the lattice enthalpy will be high.

23. General Properties of Ionic Compounds:Ionic compounds usually exist in the form of crystalline solids.

Ionic compounds have high melting and boiling points.

Ionic compounds are generally soluble in water and other polar solvents

having high dielectric constants.


5/12


6/12


6

structures, then the actual structure is said to be a resonance hybrid of these

structure.

32. Polarity of Bonds: In reality no bond is completely covalent or completely ionic.

33. Non-polar covalent bond: When a covalent bond is formed between two similar

atoms, the shared pair of electrons is equally attracted by the two atoms and

is placed exactly in between identical nuclei. Such a bond is called non-polar

covalent bond

34. Molecules having two oppositely charged poles are called polar molecules and

the bond is said to be polar covalent bond. Greater the difference in the

electro-negativity of the atoms forming the bond, greater will be the charge

separation and hence greater will be the polarity of the molecule.

35. Dipole moment is defined as the product of the magnitude of the charge and

the distance of separation between the charges.

Dipole moment ( ) = charge ( q ) x distance of separation ( d )

36. Partial Covalent Character in Ionic Bonds: When two oppositely charge ions A +

and B - are brought together; the positive ion attracts the outermost electrons

of the negative ion. This results in distortion of electron clouds around theanion towards the cation. This distortion of electron cloud of the negative ion

by the positive ion is called polarization.

37. Tendency of cation to polarize and polarisability of anion are summarized as

Fajans rules:

a. Smaller the size of the cation, greater is its polarizing power.

b. Polarisation increases with increase in size of anion. This is because the

electron cloud on the bigger anion will be held less firmly by its nucleusand, therefore, would be more easily deformed towards the cation.

c. Larger the charge on cation greater is polarizing power and larger the

charge on anion greater is its tendency to get polarized.


7/12


7

38. Valence Shell Electron Pair Repulsion (VSEPR) Theory:

Since Lewis symbols were unable to explain shapes of certain molecules,

VSEPR theory was introduced .The basic idea of this theory is that bonded

atoms in a molecule adopt that particular arrangement in space around the

central atom which keeps them on the average as far apart as possible.

39. Geometry and shapes of molecules in which central atom has no lone pair of

electrons:

Number of

electron pairs

Arrangement of

electron pairs

Molecular

geometryExamples

2 Linear

BeCl 2 ,HgCl 2

3 Trigonal planar BF 3

4 Tetrahedral CH 4 ,NH 4 +


8/12


8

5Trigonal

bipyramidalPCl 5

6 Octahedral SF 6

Shapes of simple molecules/Ions with central ions having one or more lone

pairs of electrons:

Molecule

type

No. of

bonding

pairs

No. of

lone pairs

Arrangement of

electron pairsShape Example

AB2E 2 1 Bent SO 2 ,O 3

AB3E 3 1Trigonal

pyramidalNH 3


9/12


10/12


10

41. Valence Bond Theory:

A discussion of valence bond theory is based on the knowledge of atomic

orbitals, electronic configuration of elements, overlap criteria of atomic

orbitals and principles of variation and superposition.

Orbital Overlap Concept of Covalent Bond: When two atoms approach

each other, partial merger of two bonding orbitals, known as overlapping

of the orbitals occurs.

Depending upon the type of overlapping, the covalent bonds may be

divided as sigma () bond and Pi ( ) bond.

Sigma () bond: This type of covalent bond is formed by the end to end

(hand on) overlapping of bonding orbitals along the inter-nuclear axis.

The overlap is known as head on overlap or axial overlap. The sigma bond

is formed by any one of the following types of combinations of atomic

orbitals. Sigma () bond can be formed by s s overlapping, s p

overlapping, p p Overlapping etc.

Pi ( ) Bond : This type of covalent bond is formed by the sidewise overlap

of the half- filled atomic orbitals of bonding atoms. Such an overlap is

known as sidewise or lateral overlap.

42. Hybridization:

In order to explain characteristic geometrical shapes of polyatomic

molecules concept of hybridization is used.

The process of intermixing of the orbitals of slightly different energies so

as to redistribute their energies resulting in the formation of new set of

orbitals of equivalent energies and shape.

43. Atomic orbitals used in different types of hybridization.

Shapes of molecules/ions Hybridisation type Atomic orbitals Examples

Linear sp one s + one p BeCl 2


11/12


11

Trigonal planar sp 2 one s + two p BCl 3

Tetrahedral sp 3 one s + three p CH 4 ,NH 3

Square planar dsp 2 one d +one s +two p [Ni(CN) 4 ]2-

[Pt(Cl) 4 ] 2-

Trigonalbipyramidal sp

3d one s+ three p + one d PF 5 ,PCl 5

Square pyramidal sp 3d 2 one s +three p +two d BrF 5

Octahedral sp3d 2

d 2sp 3 one s +three p +two dtwo d + one s +three p

SF 6 ,[CrF 6] 3- ,[Co(NH 3 ) 6] 3+

44. Molecular Orbital Theory (MOT):

Basic idea of MOT is that atomic orbitals of individual atoms combine to

form molecular orbitals. Electrons in molecule are present in the

molecular orbitals which are associated with several nuclei.

The molecular orbital formed by the addition of atomic orbitals is called

the bonding molecular orbital ( ).

The molecular orbital formed by the subtraction of atomic orbital is called

anti-bonding molecular orbital ( *).

The sigma ( ) molecular orbitals are symmetrical around the bond-axis

while pi ( ) molecular orbitals are not symmetrical.

Sequence of energy levels of molecular orbitals changes for diatomic

molecules like Li 2 , Be 2 , B 2 , C 2 , N 2 is 1s < *1s < 2s < *2s < ( 2p x = 2p y )


12/12


12

Bond order (b.o.) is defined as one half the difference between the number

of electrons present in the bonding and the anti-bonding orbitals.

45. Hydrogen Bonding:

The attractive force which binds hydrogen atom of one molecule withelectronegative atom like F, O or N of another molecule is known as

hydrogen bond or hydrogen bonding.

Magnitude of hydrogen bonding is maximum in solid state and least in

gaseous state.

Intermolecular hydrogen bond is formed between two different molecules

of same or different substances.

Intramolecular hydrogen bond is formed between the hydrogen atom andhighly electronegative like O, F or N present in the same molecule.c

Date post:	08-Apr-2018
Category:	Documents
Upload:	dheeraj-joshi
View:	222 times
Download:	0 times

XI Chem Ch4 ChemicalBonding&MolecularStructure Concept

Documents